1. The Halogens
Chapter 40

2. Position in Periodic TableF Cl Br I At
Gp VII

3. General information Halogen Atomic radius (nm) Ionic radius (nm)
Melting Point (oC)
Boiling Point (oC)
Density (g cm-3)
Abundance on Earth (%)
Fluorine
0.072
0.133
-220
-188
1.11
0.062
Chlorine
0.099
0.181
-101
-34.7
1.56
0.013
Bromine
0.114
0.195
-7.2
58.8
3.12
0.0003
Iodine
0.216
114
184
4.93
Astatine -
302
380
Trace

4. Characteristic properties
Very reactive non-metals
High electronegativity
High electron affinity
Bonding and Oxidation states
Colour

5. Electronegativity Element EN value F 4.0 Cl 3.0 Br 2.8 I 2.5
A measure of the attracting ability of bonding electrons.
Due to their large effective nuclear charge,
halogens are the most electronegative element
in the periodic table.

6. Electron Affinity Element EA kJ/mol F -348 Cl -364 Br -342 I -285
Generally high electron
affinity. All are exothermic.
X(g) + e-  X-(aq)

7. Bonding and Oxidation State
ns2np5
Ionic or covalent bond with oxidation state –1 or +1 (F shows –1 only)
Except F, all other halogens can expand their octet by using the low-lying, vacant d-orbitals to form bonding. Their oxidation states range from –1 to +7.

8. Bonding and Oxidation State  ns np nd
(+3 states)
HClO2, ClO2-   ns np nd
(0, +1, -1 states)
Cl2, Cl-, OCl2

9. Bonding and Oxidation State ns np nd
(+7 state)
Cl2O7, HIO4   ns np nd
(+5 state)
HClO3, HBrO3, I2O5

10. Colour Element Colour In Standard State In Water
In 1,1,1-trichloroethane F
Pale yellow Cl
Greenish yellow
Yellow Br
Reddish brown
Orange I
Violet black
Brown
Violet

11. Colour
Iodine
White light

12. Check point 40-1
E2
Fluorine
E1
Iodine

13. Variation in properties of Halogens
F Cl Br I At
400
-300
(Temp. oC)
b.p.
m.p.
m.p./b.p. increase as Mr increase
because Van der Waals’ Force increase
with Mr

14. Electronegativity 4.0 4 F form ionic compound with –1 oxidation state.
F Cl Br I 4 3 2 1
4.0
3.0
2.8
2.5
F form ionic compound
with –1 oxidation state.
I shows highly positive
oxidation states with
more electronegative
element.
e.g. KIO4

15. Electron affinity EA decrease from Cl to I due to increasing
atomic size which lowers
the nuclear attraction to
the added electron.
F has exceptional lower
EA because its 2nd shell
is already crowded with
7 e-, the additional e-
will experienced a greater
e-e repulsion than other
halogens.
F Cl Br I
280
320
340

16. Variation in Chemical Properties of elements
All are good oxidizing agents. (F2>Cl2>Br2>I2)
½ X2(std.state)  X-(aq)
X(g)
X-(g)
Hatom H
HEA
Hhyd H
Hatom
HEA
Hhyd = +

17. Oxidizing power Hatom HEA Hhyd H Oxidizing strength Element F 79.1
-348
-506
-774.9 Cl
121
-364
-607 Br
112
-342
-335
-565 I
107
-314
-293
-500
Hatom
HEA
Hhyd H
Oxidizing
strength

18. Oxidizing properties Oxidation of metals
Standard electrode potential (V)
Cl2+e-Cl-
1.36
Br2+e-Br-
1.07
I2+e-I-
0.54
Fe3++e-Fe2+
0.77
Oxidation of metals
a. F2 oxidizes all metals including Au and Ag.
b. Cl2 oxidizes common metals such as Na and Cu.
Reactions with Fe2+(aq)
2Fe2+ + Cl22Fe3+ + 2Cl-
2Fe2+ + Br22Fe3+ + 2Br-
2Fe3+ + 2I-2Fe2+ + I2

19. Oxidizing properties 3. Reaction with phosphorus 2P + 3X2  2PX3
F2 and Cl2 form PX5 as they have high oxidizing
Strength. They combine with P to exhibit its maximum
oxidation state.
Br2 and I2 form PX3 only.

20. Reaction with water Fluorine oxidizes water to form HF and O2
2F2 + 2H2O  2HF + O2
Chlorine undergoes disproportionation(self oxidation
and reduction) to form HCl and HOCl.
Cl2 + H2O  HCl + HOCl

21. Reaction with water A mixture of Cl2(aq), HCl(aq) and HOCl(aq)
is called chlorine water.
OCl- , chloric(I) ion
Strong oxidizing agent with bleaching property
OCl- + dye (coloured)  Cl- + dye (colourless)
Unstable to heat and light
2OCl-  O2 + 2Cl-

22. Reaction with water Br2 is only slightly soluble in water.
Br2 (aq) + H2O (l)  HBr(aq) + HOBr(aq)
OBr-(aq) is unstable and with bleaching property:
2OBr-  O2 + 2Br-
2. OBr- + dye(coloured)  Br- + dye(colourless)

23. Reaction with water I2 does not react with water and only very
slightly soluble in water.
I2 is more soluble in ethanol or in KI(aq).
I2(aq) + I-(aq)  I3-(aq)

24. Reaction with alkalis All halogens react with aqueous alkalis and
disproportionate in it. The products depend on
temperature
and
concentration of the alkalis.

25. Reaction with alkalis Cold dil.NaOH 2F2 + 2NaOH  2NaF + OF2 + H2O
Cl2 + 2NaOH  NaCl + NaOCl + H2O
Br2 + 2NaOH  NaBr + NaOBr + H2O
3NaOBr  2NaBr + NaBrO3
(3Br2 + 6NaOH  5NaBr + NaBrO3 + 3H2O
I2 + 2NaOH  NaI + NaOI + H2O
3NaOI  2NaI+ NaIO3
(3I2 + 6NaOH  5NaI + NaIO3 + 3H2O

26. Reaction with alkalis More concentrated NaOH
2F2 + 4NaOH  4NaF + O2 + 2H2O
Hot concentrated NaOH
Cl2 + 2NaOH  NaCl + NaCl + H2O
3NaOCl  2NaCl + NaClO3
(3Cl2 + 6NaOH  5NaCl + NaClO3 + 3H2O

27. Reaction with alkalis The reverse reaction of I2 and KOH is used
to prepare standard iodine solution for iodometric
titration. Know amount of iodine is generated by
dissolving known quantity of KIO3 in excess KI
and H+.
5KI + KIO3 + 3H+  3I2 + 6K+ + 3H2O

28. Variation in properties of halides
Halogen Displacement reactions
Oxidizing strength : F2 >Cl2 >Br2 >I2
Cl2 + 2Br-  2Cl- + Br2
Cl2 + 2I-  2Cl- + I2
Br2 + 2I-  2Br- + I2

29. Halide + c. H2SO4 NaCl + H2SO4  NaHSO4 + HCl
2NaCl + H2SO4  Na2SO4 + HCl (500oC)
2NaBr + 2H2SO4  2HBr + Na2SO4
2HI + H2SO4  SO2 + Br2 + 2H2O
2NaI + 2H2SO4  2HI + Na2SO4
8HI + H2SO4  H2S + 4I2 + 4H2O
Relative reducing power: HCl < HBr < HI

30. Halide + c. H3PO4 3NaCl + H3PO4  Na3PO4 + 3HCl
3NaBr + H3PO4  Na3PO4 + 3HBr
3NaI + H3PO4  Na3PO4 + 3HI
H3PO4 is NOT a strong oxidizing agent.
It reacts with halides to form HX.

31. Halides + Ag+(aq) Ag+(aq) + Cl-(aq)  AgCl(s) , white precipitate
Ag+(aq) + Br-(aq)  AgBr(s) , pale yellow precipitate
Ag+(aq) + I-(aq)  AgI(s) , yellow precipitate

32. Halides + Ag+(aq) Ion Action of AgNO3= Effect of Standing in sunlight
Effect of adding excess aq. NH3
Cl-
White ppt.
Turns grey
White ppt. dissolves
Br-
Pale yellow ppt.
Turns yellowish grey
Pale yellow ppt. slightly dissolves I-
Yellow ppt.
Remains yellow
Yellow ppt. does not dissolve

33. Relative acidity of HX Hydrogen Halide Dissociation constant
Degree of dissociation in 0.1M solution HF
7x10-4
0.085
HCl
1x107
0.920
HBr
1x109
0.930 HI
1x1011
0.950

34. Relative Acidity of HX H HX(aq) H+(aq) + X-(aq) H5 H6 H1
H+(g) X-(g)
H3
H4
H2
HX(g)
H (g) X (g)

35. Relative Acidity of HX H1 H2 H3 H4 H5 H6 H TS GHF 48
566
1311
-333
-1091
-515
-14
-29 15
HCl 18
431
-348
-381
-60
-13
-47
HBr 21
366
-324
-347
-64 -4 HI 23
299
-295
-305
-58 4
-62
H1
H2
H3
H4
H5
H6 H
TS G
Explanation of weak acid strength of HF:
HF has the greatest bond dissociation
energy and exceptionally small
electron affinity. It has the least
exothermic H.
Due to formation of strong hydrogen
bond and greatest degree of hydration,
HF has the smallest decrease of TS.

36. Anomalous behaviour of HF
According to La Chatelier’s Principle, the degree
of dissociation of a weak acid increases with dilution.
HA H+ + X-
However, the acidity of conc. HF is greater than
dilute HF. Why?
In conc. HF, HF molecules form dimers H2F2 by
hydrogen bond. The acid strength of H2F2 is
stronger than HF.
H2F2  H+ + HF2-

37. Uses of Halogens and Halides
Teflon
Na2SiF6 or NaF is used to fluoridate drinking water.
Chlorine bleach, sterilization of water, PVC.
Anti-knocking agent (leaded petrol)
AgBr coating in film
Iodine tincture
AgI in coating film